1. Arrangement of Electrons in Atoms Chapter 4

2. Properties of Light Electromagnetic Radiation- which is a form of energy that exhibits wavelength behavior as it travels through space. Electromagnetic Spectrum- has γ-rays, x-rays, UV, visible light, IR, Microwaves(Radar), Radio waves (TV radio, short waves, long waves) Wavelength-is the distance between points on successive waves

3. Frequency- as the number of waves that pass a given point in a specific time (s) C= speed of light 3 X 10 8 m/s. λ = wavelength. v = frequency or sometimes (f)

4. Photoelectric Effect Quantum-is the minimum quantity of energy that can be lost or gained by an atom. Photoelectric effect-refers to the emission of electrons from a metal when light shines on the metal

5. Photoelectric Effect Plank’s equation- E =hv Plank’s constant = h =6.626 X 10 -34 j s E = energy V = is the frequency of the radiation emitted

6. Quantum theory Photon- is a particle of electromagnetic radiation having zero rest mass and carrying a quantum of energy. E photon = hv Ground state- the lowest energy state of an atom Excited state- has to do with the highest energy state and PE

7. Line Emission Spectrum Lyman series- parts a-e in the UV area of Hydrogen Balmer series- parts a-e are in visible light area of Hydrogen Paschen series- parts a-d are in infrared area of Hydrogen

8. Bohr model The Bohr Model is probably familar as the "planetary model" of the atom. Electrons arranged as planets around the nucleus.

9. Quantum Mechanics Schrödinger Wave Equation (1926) – finite # of solutions  quantized energy levels – defines probability of finding an e -

10. Quantum Mechanics Heisenberg Uncertainty Principle – Impossible to know both the velocity and position of an electron at the same time

11. Quantum Numbers Orbital (“electron cloud”) – Region in space where there is 90% probability of finding an e - Orbital

12. Quantum Numbers Four Quantum Numbers: – Specify the “address” of each electron in an atom

13. Quantum Numbers Principal Quantum Number ( n ) – Energy level – Size of the orbital – n 2 = # of orbital's in the energy level

14. Quantum Numbers 2. Angular Momentum Quantum # ( l ) – Energy sublevel – Shape of the orbital

15. s

16. p

17. d

18. f

20. zn=# of sublevels per level zn 2 =# of orbital's per level zSublevel sets: 1 s, 3 p, 5 d, 7 f

21. Quantum Numbers Orbitals combine to form a spherical shape.

22. Quantum Numbers 2s 2p z 2p y 2p x

23. Quantum Numbers. Spin Quantum Number ( m s ) – Electron spin  +½ or -½ – An orbital can hold 2 electrons that spin in opposite directions.

24. Quantum Numbers

25. Pauli Exclusion Principle – No two electrons in an atom can have the same 4 quantum numbers. – Each e - has a unique “address”:

26. Quantum Numbers Principal # Ang Momentum # Magnetic # Spin # Energy level Sublevel (s,p,d,f) Orbital Electron

27. Electron configuration Electron configuration- arrangement of electrons in an atom Aufbau principle- an electron occupies the lowest energy orbital that can receive it. Pauli exclusion principle-no two electrons in the same atom can have the same set of four quantum #’s

28. Electron Configuration Hund’s rule- orbital's of equal energy are each occupied by one electron before any orbital is occupied by a second electron and all electrons in singly occupied orbital's must have the same spin.

31. Electron Configuration Noble gas configuration- using the last noble gas as a short cut method for electronic configuration.